1. Atoms

2. Looking Ahead: Atomic Models
When light from a glowing gas is dispersed into colors by a diffracting grating, each element emits a unique set of wavelengths.
We’ll learn how these spectra lines are important clues into the nature of the atom.

3. Looking Ahead: Molecules
Specialized molecules in certain microscopic plankton can emit light. These glowing waves are caused by such bioluminescent creatures.
We’ll learn why, unlike atoms, the light given off by molecules contains a broad band of wavelengths.

4. Looking Ahead: Lasers
The intense beams of light as this laser light show result from the stimulated emission of photons by atoms.
We’ll learn that when atoms give off light, they can stimulate other atoms to emit light of exactly the same wavelength and phase.

5. Looking Back: Energy-Level Diagrams
We’ve learned that the energies of atomic-sized systems are quantized. The allowed energies of such a system can be represented in an energy-level diagram.
In this chapter, we’ll use the ideas of quantization to understand in detail the energy levels and spectra of atoms and molecules.

6. Atoms
In 1987 J. J. Thomson discovered the first subatomic particle: the electron.
The mass of the electron is much smaller than the mass of the whole atom, so he thought an atom consisted of negatively charged electrons embedded in a more massive cloud of positive charge.
This model is called the “raisin-cake model.”
Thomson’s raisin-cake model of the atom.

7. Atoms
Ernest Rutherford, a former student of Thomson, discovered that a uranium crystal can emit two different rays: Beta rays and Alpha rays.
Beta rays are high-speed electrons emitted by uranium.
Alpha rays (now called Alpha particles) consist of helium nuclei, with mass m = 6.64 × 10–27 kg, emitted at high speed from a sample.
These were the first observations that atoms have an inner structure and that some atoms were not stable and could spit out various kinds of charged particles.

8. Example 1. The speed of an alpha particle
Alpha particles are usually characterized by their kinetic energy in MeV. What is the speed of an 8.3 MeV alpha particle?
solve Recall that 1 eV is the energy acquired by an electron accelerating through a 1 V potential difference, with the conversion 1.00 eV = 1.60  1019 J. First, we convert the energy to joules:
assess This is quite fast, about 7% of the speed of light.

9. The First Nuclear Physics Experiment
Rutherford set up an experiment to probe the inside of an atom. He shot high-speed alpha particles through a thin gold foil.
If the atom’s structure were described by Thomson’s raisin-cake model, then the forces exerted on the alpha particle by the positive atomic charges should roughly cancel the forces from the negative electrons, causing the particles to experience only slight deflections.

10. The First Nuclear Physics Experiment

11. The First Nuclear Physics Experiment

12. The First Nuclear Physics Experiment
Rutherford discovered that some alpha particles were deflected at large angles.

13. The First Nuclear Physics Experiment
The discovery of large-angle scattering of alpha particles led Rutherford to envision an atom in which negative electrons orbit a small, massive, positive nucleus.
This is the nuclear model of the atom.
Most of the atom is empty space.

14. Using the Nuclear Model
The atomic number of an element indicates its position in the periodic table, is the number of orbiting electrons of a neutral atom, and is the number of units of positive charge in the nucleus.
The atomic number is represented by Z.
Hydrogen, with Z = 1, has one electron orbiting a nucleus with charge +1e. Helium, with Z = 2, has two electrons orbiting a nucleus with charge +2e.

15. Using the Nuclear Model
Orbiting electrons are very light, so a photon or a rapidly moving particle, such as another electron, can knock one of the electrons away, creating a positive ion.
Removing one electron makes a singly charged ion, qion = +e. Removing two electrons creates a doubly charged ion, q = +2e.

16. Using the Nuclear Model
The positive charge of the nucleus is associated with a positive subatomic particle called the proton.
The proton’s charge is equal but opposite in sign to the electron’s charge.
A proton is about 1800 times more massive than an electron.
Electrons are more easily transferred than protons because protons are tightly bound in the nucleus.

17. Using the Nuclear Model
Scientists realized that the sum of the masses of the protons and electrons in a single atom did not equal the mass of the atom. They knew something was still missing.
In 1932, a third subatomic particle, the neutron, was discovered.
The neutron has essentially the same mass as the proton but no electric charge. It resides in the nucleus with the protons.
Neutrons help provide the “glue” that holds the nucleus together.

18. Using the Nuclear Model
A nucleus contains Z protons plus N neutrons.
The atom has a mass number A = Z + N. It is approximately the mass of the atom.
Proton and neutron masses are ≈ 1u:
1u = 1 atomic mass unit = 1.66 × 10–27 kg

19. Using the Nuclear Model
There is a range of neutron numbers that can form a nucleus with Z protons.
Such a series of nuclei are called isotopes.
The notation for isotopes is AZ, where A is the mass number.